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Electron Configurations, Orbital Notation and Quantum Numbers

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Electron Configurations, Orbital Notation and Quantum Numbers 5 Electron Configurations, Orbital Notation and Quantum Numbers Electron Configurations, Orbital Notation and Quantum Numbers Understanding Electron Arrangement and Oxidation States Chemical properties depend on the number and arrangement of electrons in an atom. Usually, only the valence or outermost electrons are involved in chemical reactions. The electron cloud is compartmentalized. We model this compartmentalization through the use of electron configurations and orbital notations. The compartmentalization is as follows, energy levels have sublevels which have orbitals within them. We can use an apartment building as an analogy. The atom is the building, the floors of the apartment building are the energy levels, the apartments on a given floor are the orbitals and electrons reside inside the orbitals. There are two governing rules to consider when assigning electron configurations and orbital notations. Along with these rules, you must remember electrons are lazy and they hate each other, they will fill the lowest energy states first AND electrons repel each other since like charges repel. Rule 1: The Pauli Exclusion Principle In 1925, Wolfgang Pauli stated: No two electrons in an atom can have the same set of four quantum numbers. This means no atomic orbital can contain more than TWO electrons and the electrons must be of opposite spin if they are to form a pair within an orbital. Rule 2: Hunds Rule The most stable arrangement of electrons is one with the maximum number of unpaired electrons. It minimizes electron-electron repulsions and stabilizes the atom. Here is an analogy. In large families with several children, it is a luxury for each child to have their own room. There is far less fussing and fighting if siblings are not forced to share living quarters. The entire household experiences a lower, less frazzled energy state. Electrons find each other very repulsive, so they too, are in a lower energy state if each “gets their own room” or in this case orbital. Electrons will fill an orbital singly, before pairing up in order to minimize electron-electron repulsions. All of the electrons that are single occupants of orbitals have parallel (same direction) spins and are assigned an up arrow. The second electron to enter the orbital, thus forming an electron pair, is assigned a down arrow to represent opposite spin. PURPOSE In this activity you will acquire an ability to write electron configurations, orbital notations and a set of quantum numbers for electrons within elements on the periodic table. You will also be able to justify oxidation or valence states using electron configurations and orbital notations. 314 Laying the Foundation in Chemistry Electron Configurations, Orbital Notation and Quantum Numbers 5 MATERIALS Periodic Table found at the end of this activity To write electron configurations and orbital notations successfully, you must formulate a plan of attack—learn the following relationships: ELECTRON CONFIGURATIONS 1. Each main energy level has n sublevels, where n equals the number of the energy level. That means the first energy level has one sublevel, the second has two, the third has three…. 2. The sublevels are named s, p, d, f, g . and continue alphabetically. The modern periodic table does not have enough elements to necessitate use of sublevels beyond f. Why s, p, d, f? Early on in the development of this model, the names of the sublevels came from sharp, principle, diffuse and fundamental, words used in describing spectral lines of hydrogen. 3. It may be easier for you to understand this by studying the table presented below: Energy level Number of Names of sublevels sublevels 1 1 s 2 2 s, p 3 3 s, p, d 4 4 s, p d, f 5 5 s, p, d, f, g Sublevel Name s p d f Number of Orbitals 1 3 5 7 Maximum number of electrons 2 6 10 14 4. Each sublevel has increasing odd numbers of orbitals available. s = 1, p = 3, d = 5, f = 7. Each orbital can hold only two electrons and they must be of opposite spin. An s-sublevel holds 2 electrons, a p-sublevel holds 6 electrons, a d-sublevel holds 10 electrons, and an f-sublevel holds 14 electrons. Laying the Foundation in Chemistry 315 5 Electron Configurations, Orbital Notation and Quantum Numbers 5. The filling of the orbitals is related to energy. Remember, electrons are lazy, much like us! Just as you would place objects on a bottom shelf in an empty store room rather than climb a ladder to place them on a top shelf, expending more energy—electrons fill the lowest sublevel available to them. Use the diagonal rule as your map as you determine the outermost or valence electron configurations for any of the elements. Using the diagonal rule you can quickly determine the electron configuration for the outermost valence electron in sulfur. First locate sulfur on the periodic table and notice that the atomic number of sulfur is 16. That means it has 16 protons and 16 electrons in a neutral atom. The first two electrons go into the 1s sublevel and fill it, the next two go into the 2s sublevel and fill it. That leaves 12 more electrons to place. The next six go into the 2p sublevel, filling it and leaving six more. Two of them go into the 3s sublevel, filling it and the remaining four go into the 3p sublevel. The completed electron configuration looks 2 2 6 2 4 like this: 1s 2s 2p 3s 3p . 6. Complete the electron configuration portion of the table on your student answer sheet. ORBITAL NOTATION Orbital notation is a drawing of the electron configuration. It is very useful in determining electron pairing and thus predicting oxidation numbers. The orbital notation for sulfur would be represented as follows: 1 2 3 4 5 8 6 9 7 10 11 12 13 16 14 15 ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑ 1s 2s 2p 3s 3p The electrons are numbered as to the filling order. Notice electrons 5, 6, and 7 went into their own orbitals before electrons 8, 9, and 10 forced a pairing to fill the 2p sublevel. This is an application of Hund’s rule which minimizes electron-electron repulsions. The same filling order is repeated in the 3p sublevel. It’s time to get the lingo straight! 316 Laying the Foundation in Chemistry Electron Configurations, Orbital Notation and Quantum Numbers 5 Electron configurations Group the 1’s, 2’s, etc. TOGETHER and it looks like this: 2 2 4 1s 2s 2p Which element has this electron configuration? Orbital notations Use blanks to represent orbitals and arrows to represent electrons and looks like this: 1 2 3 4 5 8 6 7 ↑↓ ↑↓ ↑↓ ↑ ↑ The electrons are numbered as to the filling order. 1s 2s 2p Notice electrons 5,6,7 went into their own orbitals before electron 8 forced a pairing. This minimizes repulsion. Which element has this orbital notation? 7. Complete the orbital notation column on your student answer page. JUSTIFYING OXIDATION STATES Elements in compounds have oxidation states. These oxidation states determine their behavior in the company of other elements. Your understanding of oxidation states will become very important as you learn to write correct chemical formulas for compounds. Some elements have only one oxidation state, while others have several. In general, the representative elements, those groups or families numbered as 1 - 8A have the oxidation states listed on the periodic table below. The transition metals generally have several oxidation states possible. Learn the following; it will help you make your predictions: • Metals (found to the left of the stair-step line) lose electrons to either minimize electron-electron repulsions or eliminate their valence electrons entirely. • Nonmetals tend to gain electrons to acquire an octet of electrons. An octet means the atom has eight 2 6 valence electrons arranged as ns np where n corresponds to the main energy level. Laying the Foundation in Chemistry 317 5 Electron Configurations, Orbital Notation and Quantum Numbers 2 • Transition metals generally have an oxidation state of +2 since they lose the s that was filled just before the d-sublevel began filling. • Electrons in the d-sublevels are very similar in energy to those in the s-sublevel preceding them. This means that 3d electrons are similar in energy to 4s electrons and 4d are similar to 5s, etc. • Noble gases have an octet naturally, so they generally do not react. Let’s practice. • Sulfur has many oxidation states. Use an oribital notation to justify its most common -2 oxidation state: 2 4 Sulfur has a valence electron configuration of 3s 3p . Start by drawing its orbital notation for the outermost, valence electrons. [Ne] ↑↓ ↑↓ ↑ ↑ 3s 3p Sulfur is a nonmetal and tends to gain electrons, creating the -2 charge. Gaining two electrons gives 2 6 it an octet of 3s 3p . • Copper has two common oxidation states, +1 and +2. Justify both oxidation states: 2 9 Copper has an ending electron configuration of 4s 3d . Start by drawing its orbital notation for the outermost, valence electrons. [Ar] ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ 4s 3d Since copper is a transition metal, the +2 oxidation state comes from losing the 4s electrons leaving 0 9 4s 3d . Almost all of the transition metals lose the s sublevel and have an oxidation state of +2. [Silver is an exception and only makes an oxidation state of +1.] The +1 oxidation state for copper comes from transferring one of the s electrons to the d orbitals to fill that sublevel and then losing 0 10 the remaining s electron to form 4s 3d . 8. Complete the conclusion questions that justify oxidation states on your student answer page.
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